D'ATOMIC THEORY NOTES CONTINUED
Hehe Dwew, i made it onto ye page. now serender ye notes or WALK DA PLANK!!! ARRGGHH!!
Quantum Theory:

This theory replace Bohr because his energy levels only worked for Hydrogen
The quantum theory uses complex mathematical equations to describe waves
The model predicts quantized energy levels of electrons. It depends on the probability of finding an electron in a certain position
This categorized the seven energy levels

Orbitals:

Area where an electron can be found.
There are four orbitals: s, p, d & f.
The orbitals are filled based on the location of the electrons on the periodic table.

S-Orbitals:

Given a 3-d figure, the s-orbital would look like a sphere
It can hold a maximum of two electrons – one pair of electrons
S-orbitals will occur in all seven energy levels

P-Orbitals:

Electrons begin filling the p-orbitals in the second energy level after filling the 2s – orbital
It can hold a total of six electrons – 3 electron pairs
The shape of each of the p-orbitals look like a dumbbell

D-Orbitals:

Electrons begin filling the d-orbitals in the third energy level, after filling the 4s orbital
The orbitals will hold 10 electrons – 5 electron pairs
The 3-d shape of the d-orbital is 4 x-shapes with a life-saver around the center

F-Orbitals:

They occur in the 4th energy level after filling the 6s orbital
The f-orbitals can hold 14 electrons – 7 electron pairs
The electrons are in such a shape chaos that there is no set shape to describe them.

Electron Configurations:
A visual way to write how electrons fill the orbitals

The orbitals fill as you flow left – right and top to bottom on the periodic table: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, etc.
More Notes...
DREW LANDIS AKA TEE BEE

Separation Techniques

Density- Centrifuge and Suction

When objects have different densities they can be separated layer by layer. Water is more dense than oil and it settles on the bottom of the container. An oil layer forms on the top of the container. Oil can then be suctioned off of the layer of water.

Chromatography

Separating two or more pure, dissolved liquids based on their mass
The lighter the liquids "climb" a column to be separated

Ex: water spill on a poster and the marker blends outward

Particle Size-Filtration

Just like playing in the sand at the beach, a sifter can be used to separate particles based upon their size. The bigger particles remain in the sifter while the smaller particles fall through the holes.

Charcoal Filtration

Removes color and odor from samples by adding activated charcoal to a sample, then filtering.
The charcoal collects the extra color and odor
Think of a filter system in a fish tank- the charcoal bag does this.

Ex: ash tanks/ water purifiers (brita filters)

Solubility

Different compounds are soluble in different solutions. Mixing the compounds in the liquid will create a solution and allow one compound to be rinsed from the unknown. Water is often the liquid for this process.

Heating (Endothermic)

Melting: heating a solid sample
Vaporization: heating a liquid to boiling without collecting the vapors
Distillation: heating a liquid to boiling and collecting the vapors in a new container
Sublimation: changing from solid to gas, skipping liquid (Dry Ice- solid CO2)

Cooling (Exothermic)

Freezing: changing a liquid to a solid
Condensation: collecting vapors as they return to the liquid phase
Deposition: particles settle on the surface from a vapor, solution or mixture

Chlorination

Adding chlorine to a substance to kill bacteria

Atomic Structure

Atomic Mass Notes

Atomic Mass: The average of all of the masses of the naturally occurring isotopes of an element.

The mass number is the rounded atomic mass. It is the number of protons and neutrons found in an atom.
Isotope: elements that have a different amount of neutrons

Atomic Number: The number of protons in the nucleus of an atom

Protons, Neutrons, and Electrons:

Number of Protons: same as the atomic number
Number of Electrons: same as the atomic number and the number of protons
Number of Neutrons: found by subtracting the atomic number from the mass number

Atomic Theory Notes Continued Again

Example with Directions:

Find out how many electrons are in Manganese (Mn). 25 electrons
Flow left à right and top à bottom filling the orbitals:
1s2 , 2s2 , 2p6 , 3s2 , 3p6 , 4s2 , 3d5
Self-check – add all of the superscripts; that total should match the number of electrons

Practice; Write the Electron Configurations:
v Electron Configuration for Nitrogen (N)
v Answer: 1s2 2s2 2p3
v Electron Configuration for Bromine (Br)
v Answer: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p5
v Cerium (Ce)
v Answer: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 5d1 4f1

Quantum theory notes continued again a second time

Use of Bohr’s Model:

o While Bohr’s model only worked for the Hydrogen atom, his energy ideas are still used.
o Electrons are in energy levels and can move energy levels when the electrons become energized (quantized)

Energy Levels of an Atom:

o Bohr proposed the electrons can only reside in a n energy level. The lowest energy level is closest to the nucleus
o Ground State: when electrons are in the lowest possible energy level
o Quantum: amount of energy needed to jump energy levels (particular energies given)
o Excited State: when an electron has been quantized (given particular energy) – leads to waves of light

Trends of the Periodic Table

Atomic Radius

1. Atomic radius within a group increases as one moves vertically down the periodic table
2. Atomic radius within a period decreases as one moves horizontally right across the periodic table

Group: column on periodic table
Period: row on periodic table

Trends of the Periodic Table

Ionization Energy
The energy required for elements within a group decreases as one moves vertically down the periodic table
The energy required for elements within a period increases as one moves horizontally right across the periodic table

Shielding Effect
A decrease in the attraction of the outer electrons (valence electrons) to the positively-charged nucleus
Increases as one moves vertically down the periodic table
It remains constant as you move right across the periodic table because the electrons aren’t being added to a new energy level
As you move down the periodic table there are more electrons inside and element. This creates less of an attraction to the valence electrons. (the electrons in the lower energy levels are taking up too much of the nucleus’ attention)

Electronegativity (EN)
How strong the bonds are within the compound
Decreases as you move down the periodic table because the energy levels are growing.
Increases as you move right across the periodic table because more p+ are added to the nucleus allowing for more electrons
F is the most electronegative element at 4.0 and Francium is the least at 0.7

Electron affinity
A measure of the energy change that occurs as an electron is added to an atom
Has the same trends as electronegativity for the same reasons (what I wrote on the previous slide)

Atomic Mass (Mass Number)

The Average of all of the masses of the naturally occurring isotopes of an element.
The mass number is the rounded atomic mass. It is the number of protons and neutrons found in an atom.
Isotope: The same elements with a different amount of neutrons

Atomic Mass is Expressed in Atomic Mass Units (amu)

The mass number is the mass of both the protons and neutrons, not the total mass.
Scientists developed a unit to compare all atoms.
1 amu = 1.66 x 10-24 grams (1/12th the mass of Carbon-12)

Average Atomic Mass on the PT

When you read the mass on the PT, the units are amu
Ca = 63.55 amu
These average atomic masses are the average of the atomic masses of the isotopes occurring in nature
Amu when single atom; grams when larger amounts of materials

How to calculate the amu

Scientists us the % of existence of isotopes multiplied by the mass all totaled to get the mass.
o Ex: Cu-63 exists 69.12% of the time yielding a mass of 62.94 amu and Cu-65 exists the other 30.83% of the time with a mass of 64.93 amu. Together they create the amu of Cu.
o = (0.6917 x 62.94 amu) (0.3083 x 64.93 amu) = 63.55 amu

Valence Electrons: number in s and p orbitals in that period; skip the d/f orbitals

Same group elements all have the same number of valence electrons
Group number equals the valence electron number
Max number you can have is 8 valence electrons
o Hydrogen and Helium have exception to octet (only total 2)
Ion is an element with extra fewer electrons (unbalanced)
+ ion (cations):
o left-hand side
o Groups I, II, III, (IV)
o Gave away some electrons
o Ca+2 = Calcium Ion
- ion (anions)
o right-hand side
o Group (IV), V, VI, VII, VIII
o Taking in electrons
o O-2 = Oxide Ion

Ionic Bonds

Oppositely charged ions attract and bond to each other. The compounds become known as salts.
o Ex. The attraction of Na+1 ion to a Cl-1 ion forms the compound commonly known as table salt—Sodium Choride

Writing Ionic Formulas

Write the cation and anion from the name:
o Lithium oxide: Li+1 O-2
Switch the charges to become subscripts
o Li2O
o Reduce if like charges or common multiples (+2, -2 Becomes 1:1, +4, -2 becomes 2:1)

Working From % to Formula

We can also work backwards to determine the formula of a compound. This is called the empirical formula. It gives the lowest whole-number ratio of the elements in the compound.

Step-by-Step

When given a problem assume that you have 100g total (like %, make it easy) 25.4% = 24.5 g.
Calculate the moles of each element
Divide the moles of each element by the smallest mole number to get the ratio. (Round to the nearest 0.5)
Multiply ratio by 2 if you have a 0.5 ratio
Write the formula using the whole number ratio

Sample Problem

A compound is analyzed and found to contain 25.9% N and 74.1 % O. What is the empirical formula of the compound?
o 25.9 g N x (1 mol N/14.0 g) = 1.85 mol N
o 74.1 g O x (1 mol O/15.99 g) = 4.63 mol O
o Ratio: 1.85/1.85 : 4.63/1.85 = 1 : 2.5
o Whole Number 2 (1 : 2.5) = 2 : 5
o Formula N2O5

Hydrates: compounds that have water trapped in them

The water is trapped during a cooling process
Formula: CuSO4 x 5H2O

Naming: Copper II Sulfate pentahydrate

Sample Problems:

BaCl2 x 2H2O = Barium chlorite dihydrate
MgCO3 x 5H2O = Magneseum carbonate pentahydrate
LiClO4 x 3H2O = Lithium chlorine trihydrate

% Composition of H2O

You find the total atomic mass of the compound (all elements and water)
Take the mass of the water/total mass x 100%

Sample Problems:

BaCl2 x 2H2O
Ba: 137.3 g
Cl: 2 x 35.45 = 70.90 g
H2O: 2 x 18.01 = 36.02 g
Total: 244.2 g
% H2O = (36.02 / 244.2) x 100 = 14.75%

Emperical Formula w/ Hydrates:

Treat the compound as one thing and water as another
Complete the same step-by-problems as before
Grams à moles à mole : mole ratio à formula answer

Sample Problem: Lab
10.407 g of hydrated barium iodide is heated and 0.877 g of H2O is driven off. Assuming this is all of the water, what is it’s formula and name?

10.407 – 0.877 = 9.53gBal2

Bal2 x 2H2O Barium Iodide dihydrate

Naming Covalent Compounds

1) Elements are listed by lower group first

2) If both elements are in the same group, the lower atomic number element is the first of the bond.

3) 2nd element will end with –ide

4) Prefixes are used to tell how many atoms are in the bond (mono, bi, tri, etc.)

5) First element will only have a prefix if it is more than one

Examples:

BF3 – Boron Trifluoride
Dinitrogen pentaoxide – N2O5

Introduction

Chemical reactions occur when bonds between the outermost parts of atoms are formed or broken
Chemical reactions involve changes in matter, the making of new materials with new properties, and energy changes
Symbols represent elements, formulas describe compounds, chemical equations describe a chemical reactions

Chemical Equations

Their Job: Depict the kind of reactants and products and their relative amounts in a reaction
4 Al (s) + 3 O2 (g) à 2 AL2O3 (s)
The numbers in the front are called stoichiometric coefficients.
The letters (s), (g), and (l) are the physical states of compounds.

Parts of a Reaction Equation

Chemical equations show the conversion of reactants (the molecules shown on the left of the arrow) into products (the molecules shown on the right of the arrow).
o A “+” sign separates molecules on the same side
o The arrow is read as “yields”
o Example:
§ C + O2 à CO2
o This reads “carbon plus oxygen react to yield carbon dioxide”

Chemical Equations: Why

Because of the principle of the conservation of matter
o An equation must be balanced
o It must have the same number of atoms of the same kind on both sides

Symbols Used in Equations

Solid (s)
Liquid (l)
Gas (g)
Aqueous solution (aq) (dissolve in a solution) ex: Kool-Aid
Catalyst H2SO4 (helps make the reaction) ex: heat

Escaping Gas ()
Change of temperature (D)

Balancing Equations

When balancing a chemical reaction you may add coefficients in front of the compounds to balance the reaction, but you may NOT change the subscripts
o Changing the subscripts changes the compound. Subscripts are determined by the valence electrons (charges for ionic or sharing for covalent)
Ex: H2 + O2 electricity 2H2O

à

Steps to Balancing Equations

There are four basic steps to balancing a chemical equation.
o 1. Write the correct formula for the reactants and the products
o 2. Find the number of atoms for each element on the left side. Compare those against the number of the atoms of the same element on the right side.
o 3. Determine where to place coefficients in front of formulas so that the left side has the same number of atoms as the right side for EACH element in order to balance the equation
o 4. Check your answer to see if:
§ The numbers of atoms on both sides of the equation are now balanced
§ The coefficients are in the lowest possible whole number ratios. (reduced)

Some Suggestions to Help You

Some helpful hints for balancing equations
o Take one element at a time, working left to right except for H and O. Save H for next to last, and O until last.
o IF everything balances except for O, and there is no way to balance O with a whole number, double all the coefficients and try again. (Because O is diamotic as an element)
o (Shortcut) Polyatomic ions that appear on both sides of the equation should be balanced as independent units

Diatomic Elements:

Br2, I2, N2, Cl2, H2, O2, F2

Example: C3H8 (g) + O2 (g) à CO2 + H2O (g)

Activity Series

For single replacement
“bully strength”
o Ex: Al + Fe3N2 à AlN + Fe
o Al > Fe (reaction will happen!)
o Ex: I2 + 2NaCl à 2NaI + Cl2
o I < Cl (no reaction)
§ Look at chart for answers

Solubility Chart

Double Replacement
S: soluble à label (aq)
P: partially à label (aq)
L: insoluble à label (solid; s)
H+1 = Liquid (l)
Single element à look @ P.T.
o (NH4)3PO4 + Pb(NO3)4 à Pb3(PO4) + NH4NO3

(aq) (aq) (s) (aq)

The reaction is a go!
2 (aq) products = no reaction

Limiting and Excess Reagents

In a chemical reaction, one compound drives how long and how much a reaction can occur.
Limiting: determines how much product can be formed (stops)
Excess: some left over during the reaction process

How to do the Problem:

Convert the grams of each reactant into moles of the product
The smaller mole value tells you which starting reactant is the limiting reactant.
Use the moles of the limiting reagent to calculate the amount of product produced.

LOOK ON WRITTEN SHEET (STOIC PRACTICE)

Excess Leftover?

Select one product that you formed
Convert that product into grams of excess reactant
Subtract those grams from the starting grams = how much is left

SAMPLE ON WRITTEN PAPER

Percent Yield:

Mathematical calculation of how “good” your lab results are
% yield = actual yield x 100%

theoretical yield

(the actual yield divided by the theoretical yield)

Elements that exit at gases at 25 degrees Celsius and 1 atmosphere

Helium, Nitrogen, Oxygen, Fluorine, Chlorine, Helium, Neon, Argon, Krypton, Xenon, Radon

Physical Characteristics of Gases

Gases assume the volume and shape of their containers
Gases are the most compressible state of matter
Gases will mix evenly and completely when confined to the same container.
Gases have much lower densities than liquids and solids.

Pressure = Force / Area

Units of Pressure
o 1 pascal (Pa) = 1 N/m2
o 1 atm = 760 mmHg = 760 torr
o 1 atm = 101.325 Pa = 101.325 KPa

Boyle’s Law:

Constant temperature
Constant amount of Gas
o P1 x V1 = P2 x V2
o P a 1/V
§ Note: 1’s = starting amount; 2’s = ending amount

Example:
A sample of chlorine as occupies a volume of 946 mL at a pressure of 726 mmHg. What is the pressure of the gas (in mmHg) if the volume is reduced at a constant temperature to 154 mL?

(726)(946) = P2 154
4460 mmHg

Charles Law:

As T increases (temperature) V increases (Volume)
V a T
V = constant x T
V1 / T1 = V2 / T2
Temperature must be in Kelvin
T(K) = t (degrees C) + 273.15

Example: A sample of carbon monoxide gas occupies 3.2 L at 125 degrees C. At what temperature will the gas occupy a volume of 1.54 L if the pressure remains constant?

125 C deg = 398.15
3.2 / 398.15 = 1.54 / T2
T2 = 192 K

Dalton’s Law of Partial Pressures

V and T are constant

Consider a case in which two gases, A and B, are in a container of volume V.
PA = nA RT nA is the number of moles of A
V
PB = nB RT nB is the number of moles of B
V
PT = PA + PB XA = nA XB = nB
nA + nB nA + nB
PA = XA PT PB = XB PT X = Mole fraction (%)

Lets us break apart to individual Pi = Xi PT

A sample of natural gas contains 8.2 moles of CH4, 0.421 moles of C2H6, and 0.116 moles of C3H8. If the total pressure of the gases is 1.37 atm, what is the partial pressure of propane (C3H8)?

Pi = Xi PT PT = 1.37 atm

XPROPANE = 0.166 = 0.0132
8.24 + 0.421 + 0.116

PPROPANE = 0.0132 (1.31 atm) = 0.0181 atm

With Water
PT = P02 + PH2O

SourceURL:file://localhost/Users/landisd14/Desktop/Solutions.doc

Solution: a homogenous mixture in which the components are uniformly intermingled

Solute + Solvent = Solution

Aqueous Solution: a solution in which the solvent is water

Solubility: the amount of a substance that dissolves in a given volume of solvent at a given temperature

Likes dissolve Likes (solids and liquids)

Miscible: substances that can be dissolved within one another at any ratio amount
o Ex: NaCl, H2O
§ NaCl deionizes into the Na+ and Cl- and H2O is polar having a partial so the areas of negative charge are “likes” and can be dissolve

ú 2 strong areas of negative cause a lot of action

Immiscible: substances that cannot be dissolved into one another at any ratio
o Ex: NaCl in hexane (C6H14)
§ Hexane does not have areas of negative charge to cause any “action”
Oil and Hexane readily dissolve because they both have the same type of forces; neither have an area of strong polarity
H2O and short-Carbon alcohols have areas of polarity that will dissolve easily
Examples on Paper

Why Soap Works

Soap is a long non-polar chain with an ionic tip
Grease is non-polar
Water is polar
Pg 498

SourceURL:file://localhost/Users/landisd14/Desktop/Arrhenius%20Acids%20and%20Bases.doc
Arrhenius gave a more specific definition to acids and bases aside from their feel and taste

Acids are hydrogen-containing compounds that ionize to yield H+ in aqueous solutions (Often H3O+)
Bases are compounds that ionize to yield OH in aqueous solutions

Strengths of Acids and Bases

Strength of acids and bases depend on how well the compound ionizes in water
Strong Acids- completely ionizing in water (totally break apart creating a charged particle in the water) HCl
Weak Acids- ionize only slightly (may still have a H-ion that could be removed)

Dissociation of Acids

HA + H2O à H3O + A-
Strong Acids
Weak Acids HA + H2O à H3O+ HA

Using Formulas

Strong Acids: HCl, BHr, HI and Oxoacids (#O’s > # of protons by 2 or more: H2SO4)
Weak Acids: HF; H not bonded to O, halogen; Oxoacids where #’s of O is = to or > by 1 proton: HClO; carboxylic acids (-COOH)
Strong Bases: Group 1A, Group 2A with O-2, OH-1
Weak Bases: electron-rich N compounds lacking in OH-1

Stronger Acid or Base

HClO or HClO3 HClO3
HCl or CH3COOH HCl
NaOH or CH3NH2 NaOH

Acid Dissociation Constant

Ka = Keq * [H2O] = [H3O+] * [A-]

[HA]

Tells whether an acid is strong or weak
Stronger acids have higher Ka values, higher [H3O+]
Chart on Pg. 771
[ ] = concentration (measured in molarity)

Writing Ka Expressions:

HNO2 à [H3O+][NO2-1] / HNO2

(split the H and NO2)

Hydrogen Ions and Acidity

Inside a container of water, occasional collisions between water compounds occur with enough energy to form ions
HOH + HOH à H3O+ + OH-
H3O+ are known as hydronium ions (can be simplified to H+)
This happens in small occasions and will occur in equal concentrations of ions.

Concentration of Ions

Water acts as the neutral solution when calculating the pH of any solution
The concentration (mo/L) of the ions (H3O+ and OH-) are eah said to be 1 x 10-7 M.
The overall concentration of water is then calculated as follows
o [H3O+] x [OH-] = 1 x 10-14 M = Kw
This is called the ion-product constant for water –Kw+

Acidic Concentrations

Acidic solutions are when the [H+] is greater than the [OH-]
The [H+] will be greater than 1 x 10-7 M
The [OH-] will be less than 1 x 10-7 M
o Ex: HCl à H+ + Cl-

Basic Concentrations

Basic solutions are when the [H+] is less than the [OH-] (Are also called alkaline solutions)
The [H+] will be less than 1 x 10-7
The [OH-] will be greater than 1 x 10-7
o Ex NaOH à Na+ + OH-

Calculating the pH Value

pH values range from 0-14 where 0 is extremely acidic, 7 is neutral and 14 is extremely basic
pH = -log[H+]
(Log is a button on your calculator)
pOH = -log[OH-]
pOH + pH = 14

Calculations
1) Using [H+] to solve for pH

o pH = -log[H+]

2) Using [OH-] to solve for pOH

o pOH = -log[OH-]

3) Using [OH-] to solve for pH

o pOH = -log[OH-]
o 14 = pOH

4) Using [H+] to solve for pOH

o pH = -log[H+]
o 14 = -log[H+] + pOH

5) Using pH to solve for [H+]

o pH = -log (H+)
o 10-(H+) = pH

6) Using pOH to solve for [OH-]

o pOH = -log (OH-)
o 10-(OH-) = pOH

Drew L